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CHEMICAL BONDING
Predicting deviations from ideal bond angles
Consider the carbonyl fluoride (CF2O) molecule. What is the central atom? Enter its chemical symbol.
O
How many lone pairs are around the central atom?
What is the ideal angle between the carbon-fluorine bonds?
Compared to the ideal angle, you would expect the actual angle between the carbon-fluorine bonds to (choose one) be (choose one) about the same, bigger, smaller.
Explanation
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The Correct Answer and Explanation is:
Correct Answers:
- Central atom: C
- Number of lone pairs around the central atom: 0
- Ideal angle between the carbon-fluorine bonds: 120°
- Compared to the ideal angle, the actual angle between the carbon-fluorine bonds would be: slightly smaller
Explanation (300 words):
The compound carbonyl fluoride (CF₂O) consists of a central carbon (C) atom bonded to two fluorine (F) atoms and one oxygen (O) atom. Carbon is the central atom because it is less electronegative than oxygen and typically forms four bonds, making it the logical central element in most organic molecules.
To determine the molecular geometry and predict deviations from the ideal bond angle, we apply Valence Shell Electron Pair Repulsion (VSEPR) theory. The carbon atom in CF₂O forms three sigma bonds: two with fluorine and one with oxygen. In addition, it forms a pi bond with the oxygen (as part of the carbonyl group), making a total of three regions of electron density around the central carbon atom.
According to VSEPR theory, three regions of electron density around a central atom adopt a trigonal planar geometry, which ideally results in bond angles of 120°. However, real molecules often deviate from ideal angles due to differences in electronegativity and bond character.
In CF₂O, oxygen is more electronegative than fluorine and pulls electron density toward itself, strengthening the C=O double bond. The double bond occupies more space than a single bond due to its higher electron density. This increased repulsion pushes the two C–F bonds slightly closer together, resulting in a bond angle between the two fluorine atoms that is slightly less than 120°.
Additionally, there are no lone pairs on the central carbon, so the deviation is mainly due to differences in bonding types and electronegativity effects—not lone pair repulsion.
In summary, while the ideal angle is 120°, the actual F–C–F angle is slightly smaller, due to the electron-dense C=O double bond exerting greater repulsive forces than the single C–F bonds.Correct Answers:
- Central atom: C
- Number of lone pairs around the central atom: 0
- Ideal angle between the carbon-fluorine bonds: 120°
- Compared to the ideal angle, the actual angle between the carbon-fluorine bonds would be: slightly smaller
Explanation (300 words):
The compound carbonyl fluoride (CF₂O) consists of a central carbon (C) atom bonded to two fluorine (F) atoms and one oxygen (O) atom. Carbon is the central atom because it is less electronegative than oxygen and typically forms four bonds, making it the logical central element in most organic molecules.
To determine the molecular geometry and predict deviations from the ideal bond angle, we apply Valence Shell Electron Pair Repulsion (VSEPR) theory. The carbon atom in CF₂O forms three sigma bonds: two with fluorine and one with oxygen. In addition, it forms a pi bond with the oxygen (as part of the carbonyl group), making a total of three regions of electron density around the central carbon atom.
According to VSEPR theory, three regions of electron density around a central atom adopt a trigonal planar geometry, which ideally results in bond angles of 120°. However, real molecules often deviate from ideal angles due to differences in electronegativity and bond character.
In CF₂O, oxygen is more electronegative than fluorine and pulls electron density toward itself, strengthening the C=O double bond. The double bond occupies more space than a single bond due to its higher electron density. This increased repulsion pushes the two C–F bonds slightly closer together, resulting in a bond angle between the two fluorine atoms that is slightly less than 120°.
Additionally, there are no lone pairs on the central carbon, so the deviation is mainly due to differences in bonding types and electronegativity effects—not lone pair repulsion.
In summary, while the ideal angle is 120°, the actual F–C–F angle is slightly smaller, due to the electron-dense C=O double bond exerting greater repulsive forces than the single C–F bonds.
Correct Answers:
- Central atom: C
- Number of lone pairs around the central atom: 0
- Ideal angle between the carbon-fluorine bonds: 120°
- Compared to the ideal angle, the actual angle between the carbon-fluorine bonds would be: slightly smaller
Explanation
The compound carbonyl fluoride (CF₂O) consists of a central carbon (C) atom bonded to two fluorine (F) atoms and one oxygen (O) atom. Carbon is the central atom because it is less electronegative than oxygen and typically forms four bonds, making it the logical central element in most organic molecules.
To determine the molecular geometry and predict deviations from the ideal bond angle, we apply Valence Shell Electron Pair Repulsion (VSEPR) theory. The carbon atom in CF₂O forms three sigma bonds: two with fluorine and one with oxygen. In addition, it forms a pi bond with the oxygen (as part of the carbonyl group), making a total of three regions of electron density around the central carbon atom.
According to VSEPR theory, three regions of electron density around a central atom adopt a trigonal planar geometry, which ideally results in bond angles of 120°. However, real molecules often deviate from ideal angles due to differences in electronegativity and bond character.
In CF₂O, oxygen is more electronegative than fluorine and pulls electron density toward itself, strengthening the C=O double bond. The double bond occupies more space than a single bond due to its higher electron density. This increased repulsion pushes the two C–F bonds slightly closer together, resulting in a bond angle between the two fluorine atoms that is slightly less than 120°.
Additionally, there are no lone pairs on the central carbon, so the deviation is mainly due to differences in bonding types and electronegativity effects—not lone pair repulsion.
In summary, while the ideal angle is 120°, the actual F–C–F angle is slightly smaller, due to the electron-dense C=O double bond exerting greater repulsive forces than the single C–F bonds.Correct Answers:
- Central atom: C
- Number of lone pairs around the central atom: 0
- Ideal angle between the carbon-fluorine bonds: 120°
- Compared to the ideal angle, the actual angle between the carbon-fluorine bonds would be: slightly smaller
Explanation (300 words):
The compound carbonyl fluoride (CF₂O) consists of a central carbon (C) atom bonded to two fluorine (F) atoms and one oxygen (O) atom. Carbon is the central atom because it is less electronegative than oxygen and typically forms four bonds, making it the logical central element in most organic molecules.
To determine the molecular geometry and predict deviations from the ideal bond angle, we apply Valence Shell Electron Pair Repulsion (VSEPR) theory. The carbon atom in CF₂O forms three sigma bonds: two with fluorine and one with oxygen. In addition, it forms a pi bond with the oxygen (as part of the carbonyl group), making a total of three regions of electron density around the central carbon atom.
According to VSEPR theory, three regions of electron density around a central atom adopt a trigonal planar geometry, which ideally results in bond angles of 120°. However, real molecules often deviate from ideal angles due to differences in electronegativity and bond character.
In CF₂O, oxygen is more electronegative than fluorine and pulls electron density toward itself, strengthening the C=O double bond. The double bond occupies more space than a single bond due to its higher electron density. This increased repulsion pushes the two C–F bonds slightly closer together, resulting in a bond angle between the two fluorine atoms that is slightly less than 120°.
Additionally, there are no lone pairs on the central carbon, so the deviation is mainly due to differences in bonding types and electronegativity effects—not lone pair repulsion.
In summary, while the ideal angle is 120°, the actual F–C–F angle is slightly smaller, due to the electron-dense C=O double bond exerting greater repulsive forces than the single C–F bonds.
